2.6 Batteries

Electrochemical cells used for power generation are called batteries. Although batteries come in many different shapes and sizes there are a few basic types. You won't be required to remember details of the batteries, but some general information and features of each type is presented here.

1. Primary batteries - (dry cell batteries)

• non-rechargeable

• electrolytes are present as a paste rather than as a liquid
• general purpose battery used for flashlights, transistor radios, toys, etc.
• The basic dry cell battery consists of: zinc case as the anode (oxidation); a graphite rod is the cathode (reduction) surrounded by a moist past of either MnO2, NH4Cl, and ZnCl2 or in alkaline dry cells a KOH electrolytic paste.
• General reactions for the battery - manganese(IV) oxide-zinc cell (different batteries have different reactions - you don't need to remember any of these reactions)
 cathode $$\ce{2MnO2(s) + 2NH4+ + 2e- -> Mn2O3(s) + H2O(l) + 2NH3(aq)}$$ anode $$\ce{Zn(s) -> Zn^{2+}(aq) + 2e-}$$
• Maximum voltage 1.5V. By connecting several cells in series 90V can be achieved.
• Advantages of alkaline batteries - consistent voltage, increased capacity, longer shelf-life, and reliable operation at temperatures as low as -40°C

2. Secondary Batteries (storage batteries)

• rechargeable
• an example - lead-acid battery used in cars. Anode is grid of lead-antimony or lead-calcium alloy packed with spongy lead; Cathode is lead(IV) oxide. Electrolyte is aqueous sulfuric acid. Consists of numerous small cells connected in parallel (anode to anode; cathode to cathode).
• General reaction:
 cathode $$\ce{PbO2(s) + 4H+ + SO4^{2-}(aq) + 2e- -> PbSO4(s) + 2H2O(l) + 2NH3(aq)}$$ anode $$\ce{Pb(s) + SO4^{2-}(aq) -> PbSO4(s) + 2e-}$$

• Secondary batteries are recharged by passing a current through the battery in the opposite direction. In a car battery this occurs when the engine is running.
• Other examples include the nickel-iron alkaline battery, nickel-zinc batter, nickel-cadmium alkaline battery, silver-zinc, silver-cadmium

3. Fuel Cells

• fuel cells are electrochemical cells that convert energy of a redox combustion reaction directly into electrical energy. Requires a continuous supply of reactants and a constant removal of products.
• Cathode reactant usually air or pure oxygen; anode fuel is a gas such as hydrogen, methane, or propane. Carbon electrodes typically contain a catalyst. The electrolyte is typically KOH.
• General reaction:
 cathode $$\ce{O2(g) + 2H2O(l) + 4e- -> 4OH-(aq)}$$ anode $$\ce{2H2(g) + 4OH-(aq) -> 4H2O(l) + 4e-}$$ net $$\ce{2H2(g) + O2(g) -> 2H2O(l)}$$

• Advantages - no toxic waste products (water is the only product); very efficient energy conversion (70-80% efficient)
• Disadvantage - too expensive for large-scale use.